Biological Sciences
pH Test
pH test is a method used to measure the acidity or alkalinity of a substance. It is based on the pH scale, which ranges from 0 to 14, with 7 being neutral. In biological sciences, pH tests are commonly used to assess the acidity or alkalinity of biological samples, such as blood, urine, or soil, which can have significant implications for various biological processes.
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7 Key excerpts on "pH Test"
eBook - ePub- M. Bazin(Author)
- 2018(Publication Date)
- CRC Press(Publisher)
Chapter 6Role of pH Biological Wastewater Treatment ProcessesPrasad S. Kodukula , T. B. S. Prakasam, and Arthur C. AnthonisenTable of Contents
- I. Introduction
- II. Chemistry
- A. Acid-Base Equilibria
- B. Buffers
- C. Chemical Equilibria in Biological Systems
- 1. Carbonate
- 2. Ammonia
- 3. Sulfide
- 4. Chlorine
- III. Alteration of pH by Microbial Activity
- IV. Inhibition of Enzymatic Activity
- V. Molecular Basis of Effects of pH
- VI. pH and Biological Waste Treatment
- A. Aerobic Systems
- B. Anaerobic Systems
- C. Chlorination
- VII. Summary
- References
I. Introduction
The term pH is used to describe the concentration of hydrogen ions in aqueous and other liquid media. pH is an important factor in wastewater treatment processes because of its influence on the process performance and efficiency. It plays an important role in both chemical and biological wastewater treatment systems. In the case of chemical unit operations, such as neutralization and precipitation, the treatment objective is achieved solely by adjusting the pH of the wastewater under consideration. The reaction rates in these systems are usually much faster than in biological wastewater treatment systems. In biological waste-water treatment systems, however, although the adjustment of pH is not the primary means of achieving the treatment objective, the process efficiency is nevertheless influenced by the pH, mainly because of the effect of the hydrogen ion concentration on the metabolic activity of the responsible microogranisms.In the case of some chemical treatment processes, the pH of the influent is adjusted to the desired level by addition of appropriate chemicals. Examples of such processes are (1) precipitation of phosphates1 and heavy metals2 (2) neutralization of several industrial wastes by addition of either an acid or an alkali to the process stream, depending on the inital pH of the waste, or by mixing wastes having pH values in the acidic and alkaline range3 (3) lime stabilization of sludges to destroy bacteria and viruses4 , and (4) hydrolysis of organic wastes by raising or lowering their pH through alkali or acid treatment, respectively, to increase their biodegradability.5 In the processes of adsorption6 and coagulation/flocculation,7
eBook - PDF- J. M. M. Brown, G. G. Járos(Authors)
- 2013(Publication Date)
- Butterworth-Heinemann(Publisher)
A pH of less than 7 indicates an acidic solution, while a pH of greater than 7 indicates an alkaline solution. A pH of 7 indicates neutrality in the absolute sense. 5: 12 Elementary Medical Biochemistry When we consider the hydrogen ion concentration of body fluids, neutrality of these does not coincide with the absolute value of 7, but rather with the pH values 7,3 - 7,4 i.e. the fluids of the body are slightly alkaline. In health the pH of the body fluids is maintained within these very narrow limits. A fall in pH to a value of 7,2 repre-sents a considerable increase in the hydrogen ion concentration and one which is sufficient to cause serious disturbances of cell function. If the pH of body fluids falls to 7, the increase in hydrogen ion con-centration is such that cell function is catastrophically impaired, and coma and death result. We speak of such a condition as an aci-dosis. An acidosis is a frequent complication of diseases like diabetes mellitus, in which excessive amounts of acids are produced in the body. If, on the other hand, the pH of body fluids rises above 7,6, the results can be equally catastrophic and a state of alkalosis arises. The dissociation of substances in aqueous solutions A great variety of substances dissociate in aqueous solution to yield cations and anions. Since such solutions conduct electricity, such sub-stances are referred to as electrolytes. We have encountered the fol-lowing examples so far: NaCl ^ Na + + CI K B r ^ ± K + + Br NH 4 C1 ^ NH 4 + + CI These salts, sodium chloride, potassium bromide and ammonium chloride, dissociate fully in aqueous solution and are consequently referred to as strong electrolytes. Carbonic acid can dissociate yielding hydrogen ions and bicar-bonate ions according to the equation H 2 C O 3 ^ : H + + HCO 3 ~ As in the case of water this compound dissociates poorly and solu-tions of carbonic acid contain relatively large amounts of undissoci-ated acid and few hydrogen or bicarbonate ions.
eBook - ePub- Tony Farine, Mark A. Foss(Authors)
- 2013(Publication Date)
- Routledge(Publisher)
When using the pH scale, a number of points must be considered. First, since it is a logarithmic scale, every change of one unit in pH represents a tenfold change in hydrogen ion concentration, a change of two units in pH a 100-fold change in hydrogen ion concentration, and so on. For this reason, the normal range of blood pH (7.35–7.45) is not as narrow as it first appears, and apparently small changes in blood pH represent large changes in hydrogen ion concentration, which you may need to report. Second, the pH scale is a negative scale – that is, a falling pH represents a rise in hydrogen ion concentration and a rising pH represents a falling hydrogen ion concentration.Pure water has a pH of 7 and an identical concentration of hydrogen ions and hydroxide ions, and therefore is referred to as neutral . If hydrogen ions are added, then [H+ ] rises and pH falls – that is, acids have a pH of less than 7. In contrast, if hydrogen ions are removed, then [H+ ] falls and pH rises – that is, bases have a pH of greater than 7.Salts as acids and basesWhen acids and bases react together, the salt (ionic compound) formed may be neutral, acidic or basic, depending on the strengths of the acid and base used in the reaction. If a strong acid is added to a strong base, or a weak acid is added to a weak base, then the resultant salt is neutral. In contrast, the reaction between a strong acid and a weak base results in the formation of an acidic salt, while the reaction between a weak acid and a strong base produces a basic salt.Acid–base balanceIn-text review Acids are substances that donate hydrogen ions during a chemical reaction.Bases are substances that accept hydrogen ions during a chemical reaction.Acids and bases are described as weak or strong, depending upon the extent of their dissociation.A solution of an acid or base can be concentrated or dilute, irrespective of whether the acid or base is strong or weak.Acids and bases react together to produce a salt and water.The concentration of hydrogen ions is described in terms of pH.We have already looked at the concept of homeostasis in Chapter 2
eBook - PDF- Saeed Farrokhpay(Author)
- 2023(Publication Date)
- Arcler Press(Publisher)
The pH of a solution is a measure of how basic or acidic it is. H+ activity in a solution is measured using this method, which is represented as a negative logarithm in the equation. The pH readings are presented on a scale ranging from 0.0 to 14.0 (neutral). Pure water contains a pH of 7.0 and is considered neutral; water having a pH < 7.0 is considered acidic, while water having a pH >7.0 is considered basic or alkaline. Conditions Monitoring pH and Alkalinity of Water 89 containing pH levels ranging from around 6.5 to 8.5 are preferred by most estuary species (Bae et al., 2020; Boyd, 2020). pH values are expressed on a logarithmic scale, which means that for every one-unit change in pH, acidity or alkalinity increases or decreases by a factor of ten; for example, a pH of 5.0 is 10 times more acidic than a pH of 6.0 and 100 times more acidic than a pH of 7.0. The pH of a solution is 7 when the hydroxyl and hydrogen ions are mixed in sufficient amounts (this is known as the neutral point). 3.3. THE ROLE OF PH IN THE ESTUARINE ECOSYSTEM It is possible for plants and animals to change the pH of water by photosynthesis and respiration, as well as by the minerals dispersed in the dust, aerosols, and water from the human-made pollutants and air. Human activities that result in substantial, long-term acidification of a waterbody or short-term variations in pH are extremely detrimental to the environment. For example, algal blooms, which are frequently triggered by an excess of nutrients, may affect pH levels to change substantially over a short period of time, putting a significant amount of stress on nearby species. It is possible that acid precipitation in the higher freshwater goes of an estuary will reduce the existing rate of eggs laid by breeding fish in that area (Figure 3.1) (Jarvis et al., 2006; Boyd et al., 2011). Figure 3.1. Estuarine ecosystem. Source: https://unacademy.com/lesson/estuarine-ecosystem/Y1FL6LFP.
eBook - PDF- Douglas P. Heller, Carl H. Snyder(Authors)
- 2015(Publication Date)
- Wiley(Publisher)
246 CHAPTER 8 Acids and Bases ª 4BCJOF ,BQQFMJ4UPDLQIPUP 5JN 0 8BMLFS 5JN 0 8BMLFS a. A pH meter very accurately measures hydrogen ion concentration. Here a meter records the pH of a sample of yogurt. c. Certain plants, such as beets, blueberries, cherries, and red cabbage, contain pigments that act as acid-base indicators. The figure shows the color of red cabbage extract at various pH levels. b. “Universal” test strips are saturated with combinations of acid-base indicators that turn various colors as the pH changes. Although universal strips are more sensitive than litmus paper (which simply indicates acidity or basicity), they are less accurate than pH meters. Cl − ion and one H 3 O + ion (see Figure 8.7). For example, dissolving 0.01 mole of pure HCl in 1 liter of water gives a solution with a hydronium ion concentration of 0.01 M. Between these two extremes—pure water, which ion- izes to form hydronium atoms only in the most marginal sense, and HCl, which completely ionizes in water to form hydronium atoms—we can find many examples of acids that occupy a middle ground. For example, acetic acid (the acid in vinegar) ionizes re- versibly and therefore only partial- ly in water (Figure 8.14). Acids such as acetic acid, which ionize only partially, are called weak acids. Acids such as hydrochloric acid, which ionize completely, are called strong acids. Usually pH is measured with either a pH meter or a strip of test paper. For more precise measurements, the pH meter is the instrument of choice. Most common pH meters provide values accurate to about 0.01 pH unit. Color test strips, on the other hand, give only approxi- mate values, although they are fast, convenient, and inex- pensive. We take a further look into how we measure pH in Figure 8.13. Strong Versus Weak Acids We have seen that the very slight ionization of water gen- erates a very small concentration of hydronium ions in pure water.
eBook - PDF- Charlotte W. Pratt, Kathleen Cornely(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
The pH of a solution can be altered The pH of a sample of water can be changed by adding a substance that affects the exist- ing balance between [H + ] and [OH − ]. Adding an acid increases the concentration of [H + ] and decreases the pH; adding a base has the opposite effect. Biochemists define an acid as a substance that can donate a proton and a base as a substance that can accept a proton. For example, adding hydrochloric acid (HCl) to a sample of water increases the hydrogen ion concentration ([H + ] or [H 3 O + ]) because the HCl donates a proton to water: HCl + H 2 O → H 3 O + + Cl − Note that in this reaction, H 2 O acts as a base that accepts a proton from the added acid. Similarly, adding the base sodium hydroxide (NaOH) increases the pH (decreases [H + ]) by introducing hydroxide ions that can recombine with existing hydrogen ions: NaOH + H 3 O + → Na + + 2 H 2 O TABLE 2.3 pH Values of Some Biological Fluids Fluid pH Pancreatic juice 7.8–8.0 Blood 7.4 Saliva 6.4–7.0 Urine 5.0–8.0 Gastric juice 1.5–3.0 Box 2.C Atmospheric CO 2 and Ocean Acidification The human-generated increase in atmospheric carbon dioxide that is contributing to global warming is also impacting the chem- istry of the world’s oceans. Atmospheric CO 2 dissolves in water and reacts with it to generate carbonic acid. The acid immediately dissociates to form protons (H + ) and bicarbonate (HCO 3 − ): CO 2 + H 2 O ⇌ H 2 CO 3 ⇌ H + + HCO 3 − The addition of hydrogen ions from CO 2 -derived carbonic acid therefore leads to a decrease in the pH. Currently, the earth’s oceans are slightly basic, with a pH of approximately 8.0. It has been estimated that over the next 100 years, the ocean pH will drop to about 7.8. Although the oceans act as a CO 2 “sink” that helps mitigate the increase in atmospheric CO 2 , the increase in acidity in the marine environment represents an enormous chal- lenge to organisms that must adapt to the new conditions.
eBook - ePub- Odd-Ivar Lekang(Author)
- 2019(Publication Date)
- Wiley-Blackwell(Publisher)
Below this point the carbonate system will only be represented by carbonic acid. A definition of total alkalinity can also be the requirement in the number of equivalents per litre of a strong acid to titrate the water sample down to a pH of 4.5. There can still be alkalinity down to these low pH's. Alkalinity can also be expressed as P alkalinity, meaning alkalinity above phenolphthalein indicator with endpoint pH around 8.2–8.4 but this is used in aquaculture. 5.3.3 A buffer A buffer may be added to increase the alkalinity in water when it is low. A good buffer keeps the pH constant at the required pH. For aquaculture purpose, buffers that keep pH stable at around 7 are good buffers. What is happening is that the buffer builds up alkalinity and the alkalinity is typically used/consumed in a narrow pH interval for a good buffer. Outside this interval, the pH will drop fast. The buffering capacity is strongest when the pH is close to the p K (equilibrium constant) for the media, acid buffer system. As said the typical natural buffer in water is made of the carbonate system. Here the carbonic acid–bicarbonate equilibrium has a p K value around 6.3 and the bicarbonate–carbonate equilibrium has a p K value of 10.3. This means that the titration curve flattens out in those areas (Fig. 5.4)
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